Additional physical properties of metals and important exceptions: (9) Good conductors of heat and electricity: Free electrons in metals allow easy conduction. Silver is the best conductor of electricity, followed by copper. Exception: Bismuth is a poor conductor of heat. Lead is a poor conductor compared to other metals. (10) Colour: Most metals are silver-grey or white (iron, zinc, aluminium). Exception: Gold is yellow, copper is reddish-brown. (11) Amalgam: An alloy of mercury with another metal is called an amalgam. Used in dentistry for filling cavities. Key metals and their uses: Iron (Fe) — construction, machinery. Aluminium (Al) — aircraft, packaging, utensils (light weight). Copper (Cu) — electrical wires, coins. Gold/Silver — jewellery, electronics. Zinc (Zn) — galvanisation. Mercury (Hg) — thermometers. Titanium — medical implants, aircraft. Important: Metals have 1–3 electrons in their outermost shell; they tend to lose electrons and form positive ions (cations) — hence they are electropositive elements.

Metals are elements that are generally hard, shiny, malleable, ductile, and good conductors of heat and electricity. Key physical properties: (1) Lustre: Metals have a shiny surface (metallic lustre). Gold, silver, copper have characteristic colours. (2) Hardness: Most metals are hard. Exceptions: Na and K are soft (can be cut with a knife). (3) Malleability: Metals can be beaten into thin sheets. Gold is the most malleable metal — 1g Au can be beaten into 1 sq. metre sheet. (4) Ductility: Metals can be drawn into wires. Gold and silver are most ductile — 1g Au can be drawn into 2 km of wire. (5) Tensile strength: Metals can bear large loads without breaking. (6) Sonorous: Metals produce a ringing sound when struck — used for bells. Exceptions (non-sonorous): Na, K. (7) High melting and boiling points: Most metals have high MP/BP. Exceptions: Mercury (liquid at room temp, MP = –39°C), Gallium (MP = 30°C), Caesium (MP = 29°C). (8) High density: Most metals are heavy. Exception: Na, K, Li float on water (low density).

Non-metals are elements that generally lack the properties of metals. They are non-lustrous, brittle, and poor conductors. Key physical properties: (1) Lustre: Non-metals are generally dull/non-lustrous. Exception: Iodine has a shiny appearance. (2) Hardness: Most non-metals are soft. Exception: Diamond (allotrope of carbon) is the hardest natural substance. (3) Not malleable or ductile: Non-metals are brittle — they break into pieces when beaten. (4) Non-sonorous: Do not produce ringing sound. (5) Low melting and boiling points: Most non-metals have low MP/BP. Exception: Carbon (as diamond/graphite) has very high MP. (6) Low density: Most non-metals are lightweight. (7) Poor conductors of heat and electricity. Exception: Graphite (allotrope of C) is a good conductor of electricity — used in electrodes. (8) Non-metals exist in all three states at room temperature: Solids: C, P, S, I₂. Liquid: Bromine (Br₂). Gases: O₂, N₂, H₂, Cl₂, F₂, noble gases. (9) Non-metals have 4–8 electrons in outermost shell; they tend to gain electrons — electronegative elements.

Non-metals can be solids, liquids, or gases. They are brittle, non-malleable, non-ductile, and poor conductors of heat and electricity (except graphite). Examples include Oxygen, Nitrogen, and Carbon.

Metals react with oxygen to form basic metal oxides. For example, 2Mg + O₂ → 2MgO. Sodium and Potassium react vigorously and are kept in kerosene. Some oxides like Al₂O₃ and ZnO are amphoteric.

Metals react with oxygen to form basic oxides (e.g., 4Na + O₂ → 2Na₂O). They react with water and acids to release hydrogen gas. Reactivity series (most to least reactive): K > Na > Ca > Mg > Al > Zn > Fe > Cu > Ag > Au.

Metals lose electrons to form positive ions (cations). Non-metals gain electrons to form negative ions (anions). The electrostatic attraction between them forms an ionic bond. Example: Na donates an electron to Cl to form NaCl.

The reactivity series arranges metals in descending order of reactivity. Potassium (K) is at the top, while Gold (Au) and Platinum (Pt) are at the bottom. Highly reactive metals displace less reactive ones from their salt solutions.

Metals are extracted from ores. Low reactivity metals (like Hg) are obtained by heating their sulphides. Medium reactivity metals (like Zn) are extracted by roasting or calcination followed by reduction with carbon. High reactivity metals (like Na, Al) are extracted by electrolysis.

Metals found in nature as ores are extracted by: (1) reduction with carbon for moderately reactive metals (ZnO + C → Zn + CO₂), (2) electrolytic reduction for highly reactive metals like Na, Al, Ca. Refining of metals is done by electrolysis.

Corrosion is the slow destruction of metals by environmental reaction, e.g., rusting of iron. It can be prevented by painting, oiling, greasing, galvanization (coating with zinc), or making alloys like stainless steel.

Corrosion is the deterioration of metals due to reaction with environment. Prevention methods: painting, galvanisation (coating with zinc), alloying, electroplating, and using oil/grease. Rusting of iron requires both oxygen and water.

Metals react with water to form a metal oxide or hydroxide and hydrogen gas. Potassium and sodium react violently with cold water (2Na + 2H₂O → 2NaOH + H₂); the heat evolved ignites hydrogen, so they are kept in kerosene. Calcium reacts less violently and floats because hydrogen bubbles stick to it. Magnesium reacts with hot water. Aluminium, iron and zinc react with steam: 2Al + 3H₂O → Al₂O₃ + 3H₂; 3Fe + 4H₂O → Fe₃O₄ + 4H₂. Lead, copper, silver and gold do not react with water.

Metal + dilute acid → salt + hydrogen: Mg + 2HCl → MgCl₂ + H₂; Zn + H₂SO₄ → ZnSO₄ + H₂. Copper, silver and gold do not react with dilute HCl. Hydrogen is not evolved with nitric acid, because HNO₃ is a strong oxidising agent that oxidises the hydrogen to water; only Mg and Mn give hydrogen with very dilute HNO₃. Aqua regia (3 parts conc. HCl : 1 part conc. HNO₃) can dissolve gold and platinum. A more reactive metal displaces a less reactive metal from its salt solution: Zn + CuSO₄ → ZnSO₄ + Cu; Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag.

Metal oxides that react with both acids and bases to give salt and water are amphoteric oxides. Examples: aluminium oxide (Al₂O₃) and zinc oxide (ZnO). Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (acts as base). Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (acts as acid; sodium aluminate). ZnO + 2HCl → ZnCl₂ + H₂O; ZnO + 2NaOH → Na₂ZnO₂ + H₂O. Most metal oxides are basic; non-metal oxides are acidic or neutral (CO, H₂O).

Minerals are naturally occurring compounds of elements found in the Earth's crust. Ores are minerals from which a metal can be extracted profitably. Gangue is the impurity (sand, soil, rocky material) present in ores. Examples of ores: bauxite (Al₂O₃·2H₂O) for aluminium, haematite (Fe₂O₃) for iron, cinnabar (HgS) for mercury, zinc blende (ZnS) for zinc, copper pyrites. All ores are minerals but not all minerals are ores. The ore is first enriched to remove gangue, then the metal is extracted and refined.

To reduce a metal from its ore, sulphide and carbonate ores are first converted to oxides. Roasting: heating sulphide ore strongly in excess of air: 2ZnS + 3O₂ → 2ZnO + 2SO₂. Calcination: heating carbonate ore strongly in limited air: ZnCO₃ → ZnO + CO₂. The zinc oxide is then reduced with carbon (coke): ZnO + C → Zn + CO. Metal oxides are easier to reduce than sulphides and carbonates. Reduction with carbon is used for moderately reactive metals such as Zn, Fe and Pb.

Low reactivity (Hg, Cu): heating alone. 2HgS + 3O₂ → 2HgO + 2SO₂ (roasting); 2HgO → 2Hg + O₂ (heating). Medium reactivity (Zn, Fe, Pb): convert to oxide and reduce with carbon, e.g. Fe₂O₃ + 3C → 2Fe + 3CO. Displacement with aluminium (thermite): Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat; used to weld railway tracks and cracked machine parts. High reactivity (K, Na, Ca, Mg, Al): electrolytic reduction of the molten chloride or oxide, with the metal collecting at the cathode: NaCl(molten) → Na (cathode) + Cl₂ (anode). Carbon cannot reduce these oxides, since they have more affinity for oxygen than carbon.

Impure metals obtained by reduction are refined; electrolytic refining is the most widely used method. For copper: the anode is a thick block of impure copper; the cathode is a thin strip of pure copper; the electrolyte is an acidified copper sulphate solution. On passing current, copper from the anode dissolves into the electrolyte and an equal amount of pure copper deposits at the cathode. Soluble impurities remain in solution; insoluble impurities settle below the anode as anode mud (containing gold and silver).

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal, prepared by melting the main metal and dissolving the other components. Examples: brass (copper + zinc), bronze (copper + tin), solder (lead + tin, low melting point, used for welding wires), stainless steel (iron + nickel + chromium; does not rust), steel (iron + carbon). Alloys have greater hardness, strength and resistance to corrosion than pure metals, and lower electrical conductivity. Pure gold (24 carat) is too soft, so 22 carat gold is alloyed with copper or silver. An amalgam is an alloy with mercury.